Fluorine - Wikipedia, the free encyclopedia Fluorine Distinguished from fluorene and fluorone . 9 This box: view • talk • edit Fluorine ( IPA : /ˈflʊərɪːn, -ɔːrɪːn/, Latin : fluere, meaning "to flow"), is the chemical element with the symbol F and atomic number 9. Atomic fluorine is univalent and is the most chemically reactive and electronegative of all the elements. In its elementally isolated (pure) form, fluorine is a poisonous , pale, yellowish brown gas, with chemical formula F2. Like other halogens , molecular fluorine is highly dangerous; it causes severe chemical burns on contact with skin. Fluorine's large electronegativity and small atomic radius gives it interesting bonding characteristics, particularly in conjunction with carbon. See covalent radius of fluorine . Contents [ edit ] Notable characteristics Pure fluorine (F2) is a corrosive pale yellow or brown [1] gas that is a powerful oxidizing agent. It is the most reactive and most electronegative of all the elements (4.0), and readily forms compounds with most other elements.

Its oxidation number is a constant, at -1. Fluorine even combines with the noble gases , krypton , xenon , and radon . Even in dark, cool conditions, fluorine reacts explosively with hydrogen . It is so reactive that metals , and even water , as well as other substances, burn with a bright flame in a jet of fluorine gas. It is far too reactive to be found in elemental form. In moist air it reacts with water to form also-dangerous hydrofluoric acid . In aqueous solution, fluorine commonly occurs as the fluoride ion F−, although HF is such a weak acid that substantial amounts of it are present in any water solution of fluoride at near neutral pH. Other forms are fluoro- complexes , such as [FeF4]−, or H2F+. Fluorides are compounds that combine fluorine with some positively charged counterpart. They often consist of crystalline ionic salts. Fluorine compounds with metals are among the most stable of salts.

[ edit ] History Fluorine in the form of fluorspar (also called fluorite ) ( calcium fluoride ) was described in 1530 by Georgius Agricola for its use as a flux [1] , which is a substance that is used to promote the fusion of metals or minerals . In 1670 Schwanhard found that glass was etched when it was exposed to fluorspar that was treated with acid . Carl Wilhelm Scheele and many later researchers, including Humphry Davy , Gay-Lussac , Antoine Lavoisier , and Louis Thenard all would experiment with hydrofluoric acid, easily obtained by treating calcium fluoride ( fluorspar ) with concentrated sulfuric acid. It was eventually realized that hydrofluoric acid contained a previously unknown element. This element was not isolated for many years after this, due to its extreme reactivity; fluorine can only be prepared from its compounds electrolytically, and then it immediately attacks any susceptible materials in the area. Finally, in 1886 , elemental fluorine was isolated by Henri Moissan after almost 74 years of continuous effort by other chemists.

[2] It was an effort which cost several researchers their health or even their lives. The derivation of elemental fluorine from hydrofluoric acid is exceptionally dangerous, killing or blinding several scientists who attempted early experiments on this halogen. These men came to be referred to as "fluorine martyrs." For Moissan, it earned him the 1906 Nobel Prize in chemistry (Moissan himself lived to be 54, and it is not clear whether his fluorine work shortened his life). The first large-scale production of fluorine was needed for the atomic bomb Manhattan project in World War II where the compound uranium hexafluoride (UF6) was needed as a gaseous carrier of uranium to separate the 235U and 238U isotopes of uranium . Today both the gaseous diffusion process and the gas centrifuge process use gaseous UF6 to produce enriched uranium for nuclear power applications. In the Manhattan Project , it was found that elemental fluorine was present whenever UF6 was, due to the spontaneous decomposition of this compound into UF4 and F2.

The corrosion problem due to the F2 was eventually solved by electrolytically coating all UF6 carrying piping with nickel metal, which resists fluorine's attack. Joints and flexible parts were made from Teflon , then a very recently-discovered fluorine-containing plastic which was not attacked by F2. Main article: fluoride poisoning Both elemental fluorine and fluoride ions are highly toxic and must be handled with great care and any contact with skin and eyes should be strictly avoided. When it is a free element, fluorine has a characteristic pungent odor that is detectable in concentrations as low as 20 nL/L. Its MAC value is 1 1 µL/L. All equipment must be passivated before exposure to fluorine. For more information consult an MSDS . Contact of exposed skin with hydrofluoric acid solutions poses one of the most extreme and insidious industrial threats—one which is exacerbated by the fact that hydrofluoric acid damages nerves in such a way as to make such burns initially painless. The hydrofluoric acid molecule is capable of rapidly migrating through lipid layers of cells which would ordinarily stop an ionized acid, and the burns are typically deep. HF may react with calcium, permanently damaging the bone.

More seriously, reaction with the body's calcium can cause cardiac arrhythmias, followed by cardiac arrest brought on by sudden chemical changes within the body. These cannot always be prevented with local or intravenous injection of calcium salts. Hydrofluoric acid spills over just 2.5% of the body's surface area, despite copious immediate washing, have been fatal (this corresponds with an area of about 9 in2 or 23 cm²). If the patient survives, hydrofluoric acid burns typically produce open wounds of an especially slow-healing nature. Elemental fluorine is a powerful oxidizer which can cause organic material, combustibles, or other flammable materials to ignite. Fluorocarbons are generally inert and nontoxic; the electronegativity of fluorine means that a nearby fluorine atom makes a carboxylic acid group very much more reactive. For example, trifluoroacetic acid is 100,000 times stronger than acetic acid . [ edit ] Preparation Elemental fluorine is prepared industrially by Moissan 's original process: electrolysis of anhydrous HF in which KHF2 has been dissolved to provide enough ions for conduction to take place.

In 1986, preparing for a conference to celebrate the 100th anniversary of the discovery of fluorine, Karl Christe discovered a purely-chemical preparation by reacting together at 150 °C solutions in anhydrous HF of K2MnF6 and of SbF5 . The reaction is: K 2 Mn F6 + 2 Sb F5 → 2 K Sb F6 + Mn F3 + F2 This is not a practical synthesis, but demonstrates that electrolysis is not essential. [ edit ] Compounds Fluorine forms a variety of very different compounds, owing to its small atomic size and covalent behavior, and on the other hand, its oxidizing ability and extreme electronegativity . For example, hydrofluoric acid is extremely dangerous, while in synthetic drugs incorporating an aromatic ring , fluorine is used to prevent toxication . The fluoride ion is basic, therefore hydrofluoric acid is a weak acid in water solution. However, water is not an inert solvent in this case: when less basic solvents such as anhydrous acetic acid are used, hydrofluoric acid is the strongest of the hydrohalogenic acids.

Also, owing to the basicity of the fluoride ion, soluble fluorides give basic water solutions. The fluoride ion is a Lewis base , and has a high affinity to certain elements such as calcium and silicon. For example, deprotection of silicon protecting groups is achieved with a fluoride. The fluoride ion is poisonous. Fluorine as a freely reacting oxidant gives the strongest oxidants known. Chlorine trifluoride , for example, can burn water and sand, both compounds of a weaker oxidant, oxygen. Fluorine compounds involving noble gases were first synthesised by Neil Bartlett in 1962 - xenon hexafluoroplatinate, XePtF6, being the first. Fluorides of krypton and radon have also been prepared. Also argon fluorohydride has been prepared, although it is only stable at cryogenic temperatures. The carbon-fluoride bond is covalent and very stable. The use of a fluorocarbon polymer, poly(tetrafluoroethene) or Teflon, is an example: it is thermostable and waterproof enough to be used in frying pans.

Organofluorines may be safely used in applications such as drugs, without the risk of release of toxic fluoride. In synthetic drugs, toxication can be prevented. For example, an aromatic ring is useful but presents a safety problem: enzymes in the body metabolize some of them into poisonous epoxides . When the para position is substituted with fluorine, the aromatic ring is protected and epoxide is no longer produced. Fluorine can often be substituted for hydrogen when it occurs in organic compounds. Through this mechanism, fluorine can have a very large number of compounds . Fluorite (CaF2) crystals This element is recovered from fluorite , cryolite , and fluorapatite . For a list of fluorine compounds, see here . [ edit ] Applications Chemical uses: Atomic fluorine and molecular fluorine are used for plasma etching in semiconductor manufacturing, flat panel display production and MEMS fabrication [3] . Xenon difluoride is also used for this last purpose. Hydrofluoric acid (chemical formula H F) is used to etch glass in light bulbs and other products.

Fluorine is indirectly used in the production of low friction plastics such as Teflon , and in halons such as Freon . Along with some of its compounds, fluorine is used in the production of pure uranium from uranium hexafluoride and in the synthesis of numerous commercial fluorochemicals, including vitally important pharmaceuticals, agrochemical compounds, lubricants, and textiles. Fluorochlorohydrocarbons are used extensively in air conditioning and in refrigeration . Chlorofluorocarbons have been banned for these applications because they contribute to ozone destruction and the ozone hole . Interestingly, since it is chlorine and bromine radicals which harm the ozone layer, not fluorine, compounds which do not have chlorine or bromine and contain only fluorine, carbon and hydrogen (called hydrofluorocarbons ), are not on the E.P.A. list of ozone-depleting substances [2] , and have been widely used as replacements for the chlorine and bromine containing fluorocarbons. Hydrofluorocarbons do have a greenhouse effect, but a small one compared with carbon dioxide and methane.

Sulfur hexafluoride is an extremely inert and nontoxic gas, very useful as an insulator in high-voltage electrical equipment. It doesn't occur in nature so is a useful tracer gas, though as an exceptionally potent greenhouse gas its use in unenclosed systems is inadvisable. Sodium hexafluoro aluminate ( cryolite ), is used in the electrolysis of aluminium. In much higher concentrations, sodium fluoride has been used as an insecticide, especially against cockroaches. Fluorides have been used in the past to help molten metal flow, hence the name. Some researchers including US space scientists in the early 1960s have studied elemental fluorine gas as a possible rocket propellant due to its exceptionally high specific impulse . The experiments failed because fluorine proved difficult to handle, and its combustion products proved extremely toxic and corrosive. Polytetrafluoroethylene , also known as the non-stick Teflon surface in baking pans. Dental and medical uses: Compounds of fluorine, including sodium fluoride (NaF), stannous fluoride (SnF2) and sodium MFP , are used in toothpaste to prevent dental cavities.

These compounds are also added to municipal water supplies, a process called water fluoridation , though a combination of health concerns and urban legends has sometimes led to controversy . Many important agents for general anesthesia such as sevoflurane , desflurane , and isoflurane are hydrofluorocarbon derivatives. Fluconazole is a triazole antifungal drug used in the treatment and prevention of superficial and systemic fungal infections. Halogens - humans, body, used, water, process, Earth, life, plants Halogens Halogens Photo by: Jedi Master The halogens are the five chemical elements that make up Group 17 on the periodic table: fluorine, chlorine, bromine, iodine, and astatine. The term halogen comes from Greek terms meaning "to produce sea salt." The halogens are all chemically active. For that reason, none occur naturally in the form of elements. However—with the exception of astatine—they are very widespread and abundant in chemical compounds. The most widely known of these compounds is sodium chloride, or common table salt. Fluorine and chlorine are gases. Bromine is one of only two liquid elements. Iodine is a solid.

Astatine is radioactive and is one of the rarest of the chemical elements. Fluorine is the most reactive of all known elements. Chemical reactivity decreases throughout this family of elements, with fluorine being the most reactive of all known elements, and chlorine, bromine, and iodine being relatively less reactive, respectively. Simple compounds of the halogens are called halides. When a halogen becomes part of a compound with one other element, its name is changed to an -ide ending; for example, a chloride. Fluorine The name fluorine comes from the name of the mineral in which the element was found, fluorspar. Fluorine was one of the last common elements to be isolated. It is so reactive that chemists searched for more than 70 years to find a way to extract the element from its compounds. Then, in 1886, French chemist Henri Moissan (1852–1907) found a way to produce fluorine by passing an electric current through a liquid mixture of potassium fluoride and hydrogen fluoride. Moissan's method is still used today, with some modifications, for the production of fluorine. Properties and uses.

Fluorine is one of the most dangerous chemicals known. It attacks the skin and throat, causing serious burns and respiratory problems at very low concentrations. It is also very reactive chemically. It attacks most chemicals vigorously at room temperature and reacts explosively with water. An indication of fluorine's reactivity is that it even forms compounds with the family of elements known as the inert gases. The inert gases include helium, neon, argon, krypton, and xenon. They get their name from the fact that they generally do not combine with any other element. However, compounds of xenon and fluorine and krypton and fluorine have been produced. They are the only known compounds of the inert gases to have been discovered. Words to Know Chemical activity: The tendency to form chemical compounds. Compound: The combination of two or more elements in a definite mass ratio. Radioactive: The tendency of an element to break down spontaneously into one or more other elements. Synthesized: Prepared by scientists in a laboratory; not a naturally occurring process. Because it is so reactive, fluorine itself has few uses. One exception is its role as an oxidizing (burning) agent in rocket fuels.

The vast majority of fluorine, however, is used to make compounds. One of the most interesting of those compounds is hydrofluoric acid. This compound has been used since the 1600s to etch glass. Perhaps the most familiar application of fluorine compounds is in toothpaste additives. Scientists have discovered that the addition of tiny amounts of fluoride in a person's diet can decrease the number of dental caries (cavities) that develop. Today, many kinds of toothpastes include stannous fluoride to improve a person's dental health. For many years, the most important group of fluorine compounds used commercially were the chlorofluorocarbons (CFCs). The CFCs were developed and used as refrigerants, blowing agents for polyurethane foam, and propellants in spray cans. At one time, more than 700 million kilograms (1.5 billion pounds) of CFCs were produced in a single year. In the 1980s, however, scientists found that CFCs break down in the atmosphere.

The chlorine formed as a result of this breakdown attacks the ozone layer in Earth's stratosphere (the part of Earth's atmosphere that extends 7 to 31 miles [11 to 50 kilometers] above the surface). The loss of the ozone layer is a serious problem for humans since ozone screens out radiation that causes skin cancer and other damage to plants and animals on Earth. Today, scientists are exploring the use of another class of fluorine compounds—the hydrochlorofluorocarbons, or HCFCs—as replacements for CFCs. Chlorine Chlorine was first prepared in the 1770s by Swedish chemist Carl Wilhelm Scheele (1742–1786), who thought it was a compound. It was later identified as an element by English chemist Humphry Davy (1778–1829). Davy suggested the name of chlorine for the element because of its greenish-yellow color. (The Greek word for "greenish-yellow" is chloros. ) Chlorine occurs most abundantly in sodium chloride, which is obtained from seawater and from underground deposits of rock salt formed from seas that have dried up.

To obtain chlorine, an electrical current is passed through brine, a water solution of sodium chloride. Properties and uses. Chlorine gas is toxic. It attacks the respiratory tract, causing coughing, congestion, and flu-like symptoms. In high doses, it can be fatal. For this reason, chlorine was used as a chemical weapon during World War I (1914–18). Chlorine is also very reactive, although less so than fluorine. It forms compounds with almost every other element. Among the most important of those compounds are sodium chloride (table salt), potassium chloride, hydrochloric acid, and calcium chloride. Chlorine consistently ranks among the top ten chemicals produced in the United States. Some of chlorine's uses depend on its toxic effects. For example, chlorine is now widely used as a disinfectant in municipal water systems, swimming pools, and sewage treatment plants. Many organic (carbon-containing) compounds of chlorine are used as pesticides, herbicides, and fungicides. These compounds kill unwanted insects, weeds, fungi, and other plants and animals. The use of these compounds is often associated with undesirable environmental effects, however.

DDT, for example, is a chlorine-containing compound that was once one of the most popular pesticides ever produced. But its harmful effects on fish, birds, and other animals in the environment eventually led to bans on its use in many industrialized nations. Chlorination The mention of chlorine brings summertime and swimming pools to mind for most people. Chlorine is added to pools and spas to kill bacteria in water that might otherwise cause disease. The process of adding chlorine to a swimming pool is called chlorination. Chlorination is used for other purification purposes also, as in the purification of public water supplies. Chlorination can be done in various ways. In some cases, gaseous chlorine is pumped directly into water. In other cases, a compound containing chlorine, such as sodium or calcium hypochlorite, is added to water. When that compound breaks down, chlorine and other purifying substances are released to the water. The term chlorination applies more generally to any chemical reaction in which chlorine is added to some other substance. For example, chlorine and methane gas can be reacted with each other to form a series of chlorine-containing compounds.

The best known of that series are trichloromethane (also known as chloroform; used as an anesthetic) and tetrachloromethane (also known as carbon tetrachloride; used as a solvent and a refrigerant). Chlorine is also used in the bleaching of paper, pulp, and textiles. The largest single application of the element is in the preparation of a large variety of compounds, including organic chlorides that are the starting point in the manufacture of plastics and other kinds of polymers (chemical compounds that consist of repeating structural units). One of the most important of these polymers is polyvinyl chloride (PVC), from which plastic pipe and many other plastic products are made. Another is neoprene, a synthetic form of rubber that is resistant to the effects of heat, oxidation, and oils. Neoprene is widely used in automobile parts. Bromine Bromine was discovered in 1826 by French chemist Antoine-Jérôme Balard (1802–1876). Balard chose the name bromine from the Greek word for "stink," because of its strong and disagreeable odor. Like chlorine, bromine is obtained from brine.

Chlorine gas is used to convert bromide compounds in brine to elemental bromine. Properties and uses. Bromine is a beautiful reddish-brown liquid that vaporizes (changes to a gas) easily. The vapors are irritating to the eyes and throat. The liquid is highly corrosive and can cause serious burns if spilled on the skin. Bromine is chemically less active than fluorine and chlorine but more active than iodine. Like chlorine, bromine can be used as a disinfectant. In fact, some water treatment systems have converted from chlorination to bromination as a way of purifying water. For many years, one of the most important compounds of bromine was ethylene dibromide, an additive in leaded gasolines. Since leaded gasoline has been removed from the market, this use has declined. The product in which most people are likely to encounter compounds of bromine is in photographic film. Tiny crystals of silver bromide undergo a chemical change when exposed to light. This change is responsible for the image produced when photographic film is used to take a picture. Bromine is also used to make a number of organic products that function as pesticides.

The most popular of these currently is methyl bromide, a fumigant (another word for a substance used to destroy pests). Methyl bromide is used as a spray for potatoes, tomatoes, and other agricultural crops. The halons are a group of organic compounds that contain bromine along with at least one other halogen. The halons are popular as flame retardants. However, scientists have found that, like the CFCs, they appear to cause damage to Earth's ozone layer. For that reason, their use has been largely reduced throughout the world. Iodine Iodine was discovered accidentally in 1811 by French chemist Bernard Courtois (1777–1838). Courtois was burning seaweed to collect potassium nitrate when he noticed that a beautiful violet vapor was produced. When the vapor cooled, it changed to dark, shiny, metallic-like crystals. Humphry Davy later suggested the name iodine for the element from the Greek word iodos, for "violet." As with chlorine and bromine, iodine is obtained from seawater.

It can also be produced from Chile saltpeter (sodium nitrate), in which it occurs as an impurity in the form of sodium iodate (NaIO 3 ). Properties and uses. Iodine vapor is irritating to the eyes and respiratory system. It is highly toxic if ingested. Iodine is the least active of the common halogens (not counting astatine). The human body uses iodine to make thyroxine, an important hormone (chemical messenger) produced by the thyroid gland. (The thyroid is a gland located in the neck that plays an important role in metabolism—a term used to describe processes of energy production and use by the body.) If insufficient amounts of iodine are present in the diet, a person may develop a condition known as goiter, a sometimes noticeable enlargement of the thyroid gland. Once the relationship between iodine and goiter were first discovered, manufacturers of table salt began to add iodine (in the form of sodium iodide) to their product (iodized salt). This practice has largely eliminated the problem of goiter in modern developed nations. Iodine is also used commercially in a variety of products including dyes, specialized soaps, lubricants, photographic film, medicines, and pharmaceuticals.

Astatine Astatine is generally regarded as one of the rarest naturally occurring elements. According to some estimates, no more than 44 milligrams of the element are to be found in Earth's crust. It is hardly surprising, then, that the element was first produced synthetically. In 1940, three physicists at the University of California at Berkeley—D. R. Corson, K. R. Mackenzie, and Emilio Segrè (1905–1989)—made astatine by bombarding the element bismuth with alpha particles in a cyclotron (a particle accelerator or atom-smasher). About 24 isotopes (forms) of astatine exist, all of them radioactive. The most long-lived has a half-life of 8.3 hours, meaning that half of a sample of the element disappears in 8.3 hours. Because it is so rare and has such a short half-life, astatine is one of the most poorly understood of all chemical elements. It has no practical applications at this time.

Fluorine (F) - Chemical properties, Health and Environmental effects Moissan in 1886 Fluorine Fluorine is an univalent poisonous gaseous halogen, it is pale yellow-green and it is the most chemically reactive and electronegative of all the elements. Fluorine readily forms compounds with most other elements, even with the noble gases krypton, xenon and radon. It is so reactive that glass, metals, and even water, as well as other substances, burn with a bright flame in a jet of fluorine gas. In aqueous solution, fluorine commonly occurs as the fluoride ion F-. Fluorides are compounds that combine fluoride with some positively charged counterpart. Applications Atomic fluorine and molecular fluorine are used for plasma etching in semiconductor manufacturing, flat panel display production and MEMs fabrication. Fluorine is indirectly used in the production of low friction plastics such as teflon and in halons such as freon, in the production of uranium. Fluorochlorohydrocarbons are used extensively in air conditioning and in refrigeration.

Fluorides are often added to toothpaste and, somewhat controversially, to municipal water supplies to prevent dental cavities. Fore more information visit our page on mineral water . Fluorine in the environment Annual world production of the mineral fluorite in around 4 million tonnes, and there are around 120 million tonnes of mineral reserves. The main mining areas for fluorite are China, Mexico and Western Europe. Fluorine occurs naturally in the earth's crust where it can be found in rocks, coal and clay. Fluorides are released into the air in wind-blown soil. Fluorine is the 13th most aboundant element in the Earth's crust: 950 ppm are contanined in it. Soils contain approximatively 330 ppm of fluorine, ranging from 150 to 400 ppm. Some solis can have as much as 1000 ppm and contaminated solis have been found with 3500 ppm. Hydrogen fluorides can be released into air through combustion processes in the industry. Fluorides that are found in air will eventually drop onto land or into water.

When fluorine is attached to very small particles it can remain in the air for a long period of time. In the atmosphere 0.6 ppb of fluorine are present as salt spray and organicochloride compounds. Up to 50 ppb has been recorded in city environments. Health effects of fluorine Small amounts of fluorine are naturally present in water, air, plants and animals. As a result humans are exposed to fluorine through food and drinking water and by breathing air. Fluorine can be found in any kind of food in relatively small quantities. Large quantities of fluorine can be found in tea and shellfish. Fluorine is essential for the maintenance of solidity of our bones. Fluorine can also protect us from dental decay, if it is applied through toothpaste twice a day. If fluorine is absorbed too frequently, it can cause teeth decay, osteoporosis and harm to kidneys, bones, nerves and muscles. Fluorine gas is released in the industries. This gas is very dangerous, as it can cause death at very high concentrations. At low concentrations it causes eye and nose irritations.

Flourine Testing | Elemental Testing and Analysis | Midwest Microlab Home » Flourine Testing | Elemental Testing and Analysis | Midwest Microlab » Fluorine Testing Flourine Elemental Testing and Analysis Established in 1956, Midwest Microlab Inc. is equipped to provide rapid accurate organic microchemical analysis to research and industry. Specialties include analysis for Carbon, Hydrogen, Nitrogen, direct Oxygen, and halogens including Fluorine. What is Fluorine? Fluorine (from L. fluere, meaning “to flow”), is the chemical element in the periodic table that has the symbol F and atomic number 9. Atomic fluorine is univalent and is the most chemically reactive and electronegative of all the elements. In its pure form, it is a poisonous, pale, yellow-green gas, with chemical formula F2. Like other halogens, molecular fluorine is highly dangerous; it causes severe chemical burns on contact with skin. Fluorine’s relatively large electronegativity and small atomic radius makes it have some interesting bonding characteristics. Pure fluorine (F2, since fluorine is diatomic) is a corrosive pale yellow gas that is a powerful oxidizing agent.

It is the most reactive and electronegative of all the elements, and readily forms compounds with most other elements. Fluorine even combines with the noble gases, krypton, xenon, and radon. Even in dark, cool conditions, fluorine reacts explosively with hydrogen. It is so reactive that glass, metals, and even water, as well as other substances, burn with a bright flame in a jet of fluorine gas. It is far too reactive to be found in elemental form and has such an affinity for most elements, including silicon, that it can neither be prepared nor be kept in glass vessels. In moist air it reacts with water to form the equally dangerous hydrofluoric acid. In aqueous solution, fluorine commonly occurs as the fluoride ion F-. Other forms are fluoro-complexes, such as [FeF4]-, or H2F+. Fluorides are compounds that combine fluorine with some positively charged counterpart. They often consist of ions. Fluorine compounds with metals are among the most stable of salts. Analysis of Fluorine using the ISE: The appropriate size solid sample is weighed and inserted into a micro bomb.

The bomb is then charged with the appropriate reagents to completely digest the fluorine compound. The apparatus is clamped shut and heated in a flame for 60 seconds. The contents are washed out and diluted. The readings taken with the electrode are compared against readings from two dilutions of NaF standard. Thus a standardization check is done every time fluorine is determined. The preceding is a brief description of our techniques used in fluorine determinations. With air-sensitive and/ or liquid samples digestion is done after sealing the sample in a Sn Capsule. Fluorine is a complex and labor intensive determination. It is important to note that the sample does not come into contact with glass through out the process. Aluminum and -OH are the only know interferences. Strong base is circumvented in the dilution process. With some modifications this determination is carried out according to the method of Ma and Rittner (“Modern Organic Elemental Analysis” Marcel Dekker Inc. New York, NY 1979) Please phone or e-mail us if you have specific questions about Fluorine analysis in regard to your sample.

The Parts of the Periodic Table Lr Group 8A (or VIIIA) of the periodic table are the noble gases or inert gases: helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and radon (Rn). The name comes from the fact that these elements are virtually unreactive towards other elements or compounds. They are found in trace amounts in the atmosphere (in fact, 1% of the atmosphere is argon); helium is also found in natural gas deposits. In their elemental form at room temperature, the Group 8A elements are all colorless, odorless, monatomic gases. The Group 8A elements have a full octet of eight valence electrons in their highest-energy orbitals (ns2np6), so these elements have very little tendency to gain or lose electrons to form ions, or share electrons with other elements in covalent bonds. They can be forced to form compounds with other elements, but require special conditions to do so.

Xenon, krypton compounds are known (this was first accomplished by Neil Bartlett in 1962), including XeF2, XeF4, XeF6, XeOF2, XeOF4, XeO2F2, XeO3F2, XeO2F4, XeO3, XeO4, KrF2, RnF2, among others. Helium (He, Z=2). Helium is a colorless, odorless, unreactive gas which liquefies at -268.97ï¿½C (4.18 K). The name "helium" is derived from the Greek word for the Sun, helios. Helium was originally discovered on the Sun in 1868 by Pierre Janssen: his analysis of the light emitted from the Sun's corona during a solar eclipse showed that there were lines in the Sun's spectrum that were produced by a previously unknown element.

When helium was found on Earth in the late 1890's (in association with uranium ores), it was found to have the same emission spectrum as the element previously found in the Sun, indicating that it was indeed the same element. The concentration of helium in the Earth's crust is about 8 ppb, making it the 71st most abundant element; it is found in the atmosphere at a concentration of 5 ppm (by volume). It is found in some minerals, where it is produced by radioactive elements that decay by alpha particle emission (see below). Helium is the second most abundant element in the universe (23% by mass); hydrogen and helium together make up 99% of the "normal" matter of the universe. (Of course, there's also "dark matter" and "dark energy" to worry about, but that's another story.) Hydrogen, helium, and trace amounts of lithium were produced at the beginning of the Universe in the Big Bang, and became concentrated into stars by the force of gravity. The fusion of hydrogen atoms in stars to produce helium produces huge amounts of energy; the energy in sunlight is captured by plants in photosynthesis, and drives most of the chemistry of living organisms.

The most common isotope of helium, helium-4, consists of two protons and two neutrons in the nucleus, surrounded by two electrons. An alpha particle is a type of radiation emitted by some radioactive nuclei, which consists of two protons and two neutrons, having a 2+ charge; when an alpha particle is emitted by a radioactive element, it quickly picks up two electrons from its environment, becoming a helium atom. Helium is found along with some radioactive minerals that decay by alpha particle emission, however the concentration of helium in these minerals is very low. The main source of helium is from natural gas, where it is separated out from the other components by fractional distillation. Helium liquefies at 4.18 K (-268.97ï¿½C); at 2.17 K it becomes a superfluid, in which all viscosity (resistance to flow) is lost. Some materials become superconductors when immersed in liquid helium; in these materials, there is no resistance to the flow of electricity. Superconducting magnets are used for an important medical imaging technology called Magnetic Resonance Imaging (MRI), which allows images of tissues and organs to be obtained.

A related technology is used by chemists to obtain information that can be used to deduce the structures of molecules; this technique is called Nuclear Magnetic Resonance (NMR); the two technologies are basically the same, but the term "nuclear" is avoided in the medical application because of its unpleasant associations (even though there is no nuclear radiation involved). Helium is also used in blimps, where the combination of being lighter than air and nonflammable makes it much safer than hydrogen (as demonstrated in the explosion of the Hindenburg ). It is also used in deep-sea diving, since helium is less soluble in the bloodstream than nitrogen, and reduces the risk of the diver suffering nitrogen narcosis, high pressure nervous syndrome (HPNS), or the bends. Breathing helium gas causes the voice to become higher pitched because helium is less dense than air, changing the frequency of the sound waves generated in the larynx. Neon (Ne, Z=10). Neon is a colorless, odorless, unreactive gas which liquefies at -246.1ï¿½C (27.1 K).

Neon was discovered in 1898 by Sir William Ramsay (see entry on Argon below) and his assistant Morris Travers while working with a sample of krypton. The name comes from the Greek word for "new," neos. It is found in the Earth's crust at a concentration of 70 ppt, making it the 80th most abundant element; it is found in the atmosphere at a concentration of 18 ppm (by volume). Neon which is used commercially is extracted from liquid air by fractional distillation. Neon is used primarily in neon lights; stimulating neon with electricity causes it to emit an intense red light. Other noble gases are used in "neon lights" as well, to produce different colors. Neon is also used in some underwater breathing mixtures; liquid neon is used in low-temperature cooling. Argon (Ar, Z=18). Argon is a colorless, odorless, unreactive gas which liquefies at -185.8ï¿½C (87.3 K).

It is found in the Earth's crust at a concentration of 1.2 ppm (produced by the radioactive decay of potassium-40), making it the 56th most abundant element; it is found in the atmosphere at a concentration of 0.93% (by volume). Argon which is used commercially is extracted from liquid air by fractional distillation. Even though argon makes up almost 1% of the Earth's atmosphere, it was not identified until 1894 by John William Strutt, 3rd Baron Rayleigh (Nobel Prize in Physics, 1904) and Sir William Ramsay (Nobel Prize in Chemistry, 1904). Argon had been observed before (by Henry Cavendish in 1785), but it had not been identified as an element. Because of its unreactivity, it was named for the Greek word for "idle," argos. When argon was discovered, it didn't fit with the reactivities of any other element on the periodic table, and chemists realized that there might be a whole family of previously-unrecognized elements, which we now recognize as Group 8A.

Argon is used in light bulbs because it is inert to the hot metal of the bulb filament; the filament would burn out much faster in air or even pure nitrogen. It is also used in metal refining, some "neon" lights (argon produces a blue light), some food packaging (to provide an inert atmosphere to reduce spoilage), in the tires of some luxury cars in place of air, and in lasers used in eye surgery. Argon is produced by the radioactive decay of potassium-40; the potassium-40 nucleus undergoes electron capture to produce the stable argon-40 nucleus. The ratio of potassium-40 to argon-40 in some rocks can be used to determine the age of the rock. Recently, the first argon compound, argon fluorohydride (HArF), was reported , which is stable up to 27 K. (I'm waiting for hydrogen to be replaced by boron: I think a molecule that spells out "BArF" would be worth investigating. Or not.) Krypton (Kr, Z=36).

Krypton is a colorless, odorless, unreactive gas which liquefies at -153.2ï¿½C (119.9 K). It is found in the Earth's crust at a concentration of 10 ppt, making it the 81st most abundant element; it is found in the atmosphere at a concentration of 1 ppm (by volume). Krypton which is used commercially is extracted from liquid air by fractional distillation. Krypton was discovered in 1898 by Sir William Ramsay and his assistant, Morris Travers, in a sample of argon gas extracted from the atmosphere. The name "krypton" is derived from the Greek word kryptos, meaning "hidden" (which has nothing to do with the home planet of Superman). Krypton is used in some "neon" lights, producing a violet color when electricity is passed through it. Radioactive kyrypton-85 is produced in nuclear reactors and nuclear processing plants; monitoring krypton-85 in the atmosphere allowed the United States to monitor how much nuclear material the Soviet Union was using. Xenon (Ze, Z=54).

Xenon is a colorless, odorless, unreactive gas which liquefies at -111.7ï¿½C (161.4 K). It is found in the Earth's crust at a concentration of 2 ppt, making it the 83rd most abundant element; it is found in the atmosphere at a concentration of 90 ppb (by volume). Xenon which is used commercially is extracted from liquid air by fractional distillation. Xenon was also discovered by Ramsay and Travers in 1898. The name is derived from the Greek word for "stranger," xenos. Xenon lights glow with a blue light, and also emits some low-frequency ultraviolet light. Xenon lights are used in tanning beds, biocidal lamps used in food preparation, car headlights, flash lighting on cameras, strobe lights used in high-speed photography, and has been tested for use in space travel in ion-propulsion engines (such as the experimental NASA spacecraft Deep Space 1 ). Radon (Rn, Z=86).

Radon is a colorless, odorless, radioactive gas which liquefies at -61.8ï¿½C (211.3 K). It is found in the Earth's crust in only trace amounts, and is one of the ten least abundant elements; it is found in the atmosphere at a concentration of 10-9 ppt (by volume). Radon was discovered in 1900 by Friederich Ernst Dorn while investigating "radium emanation" (first observed by Marie and Pierre Curie), a gas found in sealed vessels of radium. The name is a variation on the name of the element radium. It is produced in the radioactive decay of uranium and thorium, which in a series of steps decay into radium, which then undergoes alpha-decay to produce radon. The most common isotopes of radon are radon-220 and radon-222, which have half-lives of 55.6 seconds and 3.82 days, respectively. Although chemically unreactive, when radon is produced, it can escape from the ground in the gas form, and collect in basements, or other contained areas.

Radon is an alpha-particle emitter; while alpha particles can be blocked easily, if the radon gas is inhaled, and undergoes alpha decay in the lungs, it produces an isotope of polonium, which is a solid, as are all of its decay products, some of which have half-lives of over 20 years. Having alpha-particle emitters inside the lungs leads to a greatly increased risk of lung cancer, accounting for about 10% of all reported cases of lung cancer. References John Emsley, The Elements, 3rd edition. Oxford: Clarendon Press, 1998. John Emsley, Nature's Building Blocks: An A-Z Guide to the Elements. Oxford: Oxford University Press, 2001. David L. Heiserman, Exploring Chemical Elements and their Compounds. New York: TAB Books, 1992. Properties of Oxygen About Watch and Favorite Watch Watching this resources will notify you when proposed changes or new versions are created so you can keep track of improvements that have been made. Favorite Favoriting this resource allows you to save it in the “My Resources” tab of your account.

There, you can easily access this resource later when you’re ready to customize it or assign it to your students. Properties of Oxygen Oxygen is a highly reactive nonmetallic element; it is a strong oxidizing agent with high electronegativity and forms O2 at Standard Temperature and Pressure (STP). Learning Objective Review the properties of oxygen. Key Points At standard temperature and pressure (STP), two atoms of the element bind to form dioxygen, a colorless, odorless, tasteless diatomic gas with the formula O2. Oxygen is a member of the chalcogen group on the periodic table and is a highly reactive nonmetallic element. As such, it readily forms compounds (notably oxides) with almost all other elements. Oxygen is a strong oxidizing agent and has the second-highest electronegativity of all reactive elements, second only to fluorine. The solubility of oxygen in water is temperature-dependent; it condenses at 90.20 K and freezes at 54.36 K.

Terms ozone An allotrope of oxygen (symbol O3) having three atoms in the molecule instead of the usual two; it is a blue gas, generated from oxygen by electrical discharge; it is poisonous and highly reactive, but it protects life on Earth by absorbing solar ultraviolet radiation in the upper atmosphere. Register for FREE to remove ads and unlock more features! Learn more Full Text Oxygen is an important part of the atmosphere and is necessary to sustain terrestrial life. Because it comprises most of the mass in water, it also comprises most of the mass of living organisms. All major classes of structural molecules in living organisms, such as proteins, carbohydrates , and fats, contain oxygen, as do the major inorganic compounds that comprise animal shells, teeth, and bone. Elemental oxygen (O2) is produced by cyanobacteria, algae, and plants through the process of photosynthesis , and is used in cellular respiration by most living organisms on earth. Oxygen is toxic to obligate anaerobic organisms (organisms which need a lack of oxygen for survival), which were the dominant form of early life on Earth, until O2 began to accumulate in the atmosphere.

Chemical Properties of Oxygen At standard temperature and pressure (STP), two atoms of the element bind to form dioxygen, a colorless, odorless, tasteless diatomic gas with the formula O2. Oxygen is a member of the chalcogen group on the periodic table and is a highly reactive nonmetallic element. As such, it readily forms compounds (notably, oxides) with almost all other elements. Oxygen is a strong oxidizing agent and has the second-highest electronegativity of all reactive elements, second only to fluorine. By mass, oxygen is the third-most abundant element in the universe, after hydrogen and helium, and the most abundant element by mass in the Earth's crust, making up almost half of the crust's mass. Free oxygen is too chemically reactive to appear on Earth without the photosynthetic action of living organisms, which use the energy of sunlight to produce elemental oxygen from water. Elemental O2 only began to accumulate in the atmosphere after the evolutionary appearance of photosynthetic organisms, roughly 2.5 billion years ago. Diatomic oxygen gas currently constitutes 20.8 percent of the volume of air.

Diatomic Oxygen The two oxygen atoms in diatomic oxygen are chemically bonded to each other with a spin triplet electron configuration . This bond has a bond order of two and is often simplified in descriptions as a double bond , or as a combination of one two-electron bond and two three-electron bonds. Triplet oxygen (not to be confused with ozone , O3) is the ground state of the O2 molecule. The electron configuration of the molecule has two unpaired electrons occupying two degenerate molecular orbitals . These orbitals are classified as antibonding (weakening the bond order from three to two), so the diatomic oxygen bond is weaker than the diatomic nitrogen triple bond , in which all bonding molecular orbitals are filled, but some antibonding orbitals are not. In normal triplet form, O2 molecules are paramagnetic . This means they behave as magnets in the presence of an external magnetic field, because of the spin magnetic moments of the unpaired electrons in the molecule. Liquid oxygen is attracted to a magnet to a sufficient extent that, in laboratory demonstrations, a bridge of liquid oxygen may be supported against its own weight between the poles of a powerful magnet.

Singlet oxygen is a name given to several higher-energy species of molecular O2 in which all the electron spins are paired. It is much more reactive toward common organic molecules than is the triplet form of molecular oxygen. Physical Properties of Oxygen Oxygen is more soluble in water than nitrogen is; water contains approximately one molecule of O2 for every two molecules of N2, compared to an atmospheric ratio of approximately one to four. The solubility of oxygen in water is temperature-dependent, and about twice as much (14.6 mg/L) dissolves at 0 °C than at 20 °C (7.6 mg/L). At 25 °C and 1 standard atmosphere (101.3 kPa) of air, freshwater contains about 6.04 milliliters (mL) of oxygen per liter , whereas seawater contains about 4.95 mL per liter. At 5 °C the solubility increases to 9.0 mL (50 percent more than at 25 °C) per liter for water and 7.2 mL (45 percent more) per liter for sea water.

Oxygen condenses at 90.20 K (−182.95 °C, −297.31 °F), and freezes at 54.36 K (−218.79 °C, −361.82 °F). Both liquid and solid O2 are clear substances with a light sky-blue color caused by absorption in the red (in contrast with the blue color of the sky, which is due to Rayleigh scattering of blue light) .

Question: Which gas is the most chemically reactive to all the elements, to the extentthat water will burn in it?
Answer: FLOURINE